Thermochemistry Flashcards
1
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Types of Systems
A
- Isolated: System cannot exchange energy (heat and work) or matter with the surroundings, such as bomb calorimeter.
- Closed: System can exchange energy (heat and work) but not matter with the surroundings, such as steam radiator.
- Open: System can exchange both energy (heat and work) and matter with the surroundings, such as pot of boiling water.
2
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First Law of Thermodynamics
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- ΔU = Q - W, where ΔU is change in internal energy of system, Q is heat added to system, and W is work done by system.
- Isothermal Process: Occurs when system’s temperature is constant. Constant temperature implies that U is constant throughout process and that ΔU = 0, simplifying first law to W = Q (heat added to the system equals work done by the system). Area under Pressure vs Volume curve represents work performed by gas, as well as heat that entered the system.
- Adiabatic Process: Occurs when no heat is exchanged between the system and the environment. Thus Q = 0, simplifying first law to ΔU = -W (change in internal energy of the system is equal to work done on the system by the environment). If ΔU < 0, temperature decreases; if ΔU > 0, temperature increases.
- Isobaric Process: Occurs when the pressure of the system is constant. Does not alter first law, but does appear as a horizontal line on P-V graph.
- Isovolumetric (Isochoric) Process: Occurs when volume is unchanged. Constant volume implies that the gas neither expands or compresses and that no work is performed, simplifying first law to ΔU = Q (change in internal energy is equal to heat added to the system). Appears as vertical line on P-V graph; area under graph represents work done by the gas, which is zero.
3
Q
Standard Conditions and STP
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- Standard Conditions: 25°C (298K), 1atm pressure, 1M concentrations. Used for kinetics, equilibrium, and thermodynamics.
- Standard Enthalpy (ΔH°), Standard Entropy (ΔS°), Standard Free Energy (ΔG°) are all under standard conditions.
• Standard Temperature and Pressure (STP): 0°C (273K) and 1atm. Used for ideal gases.
4
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Phase Changes
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- Line A represents solid-liquid interface, Line B represents the liquid-gas interface, Line C represents the solid-gas interface.
- Gas phase found at high temperatures and low pressures.
- Liquid phase found at moderate temperatures and moderate pressures.
- Solid phase found at low temperatures and high pressures.
- Triple Point: Intersection of the three phase boundaries.
- Critical Point: Termination point of liquid-gas phase boundary at which there is no distinction between gas and liquid phase.
- Solid-liquid line of equilibrium extends indefinitely from triple point.
5
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Heat
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- Heat (Q): The transfer of energy from one substance to another as a result of their differences in temperature. Heat is equivalent to Enthalpy (ΔH) under constant pressure.
- Temperature is the average kinetic energy of molecules in a substance and is proportional to substance’s internal energy.
- Endothermic (ΔQ > 0) processes absorb heat.
- Exothermic (ΔQ < 0) processes release heat.
- Heat Transfer: q = mcΔT, where m is mass (g), c is specific heat of substance, and ΔT is change in temperature.
- Specific heat of H₂O (l) is 1.0 cal/(g•K) or 4.184 J/(g•K).
- Heat Capacity = mc.
- Bomb Calorimeter is constant-volume and adiabatic (no heat transfer between calorimeter and rest of universe). Heat transfer occurs only within calorimeter between system and surroundings. Constant volume means W = 0.
- So, q(system) = -q(surroundings).
- So, mcΔT(sample) + mcΔT(oxygen) = -mcΔT(water).
• Heating Curves and Phase Changes: When substance is in one phase increasing its temperature, q=mcΔT is used. When substance is undergoing phase change, temperature does not change and q=mΔH or q=nΔH must be used, where ΔH can be either Enthalpy of Fusion ΔH(fus) or Heat of Vaporization ΔH(vap).
6
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Enthalpy
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- Enthalpy (ΔH) is equivalent to heat transferred into and out of system at constant pressure.
- ΔH(rxn) = H(products) - H(reactants).
- +ΔH(rxn) implies Endothermic process.
- -ΔH(rxn) implies Exothermic process.
- Standard Enthalpy of Formation of compound ΔH°(f) is the enthalpy required to produce one mole of a compound from its elements in their standard states at 298K and 1atm.
- Standard Enthalpy of Rxn ΔH°(rxn) = ΣΔH°(f,products) - ΣΔH°(f,reactants).
- Hess’s Law: Enthalpy changes of reactions are additive; reverse reactions have same magnitudes but opposite signs, and enthalpies must be multiplied by coefficients when adding.
• Bond Dissociation Energy (Bond Enthalpy): ΔH°(rxn) = ΣΔH(bonds broken) - ΣΔH(bonds formed) = Total Energy Absorbed - Total Energy Released.
7
Q
Entropy
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- Second Law of Thermodynamics: Energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so.
- ΔS = Q/T, with units of J/(mol•K).
8
Q
Gibbs Free Energy
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- ΔG = ΔH - TΔS.
- -ΔG implies spontaneous Exergonic reaction.
- +ΔG implies nonspontaneous Endergonic reaction.
- ΔG°(rxn) = -RT ln(Keq).
- Reactions with equilibrium constants K(eq) > 1 have -ΔG and are spontaneous.
- Reactions with equilibrium constants K(eq) < 1 have +ΔG and are nonspontaneous.
- ΔG(rxn) = RT ln(Q/Keq).
- If Q < K(eq), forward reaction is spontaneous.
- If Q > K(eq), reverse reaction is spontaneous.
- If Q = K(eq), reaction is at equilibrium and ΔG = 0.
- In the presence of a catalyst, overall free energy (difference between free energy of products and reactants) will be unchanged, but activation energy will be reduced.
9
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Thermodynamics vs Kinetics
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- Rate of reaction depends on Activation Energy, not Gibbs Free Energy.
- When a reaction is thermodynamically spontaneous, it has no bearing on how fast it goes to equilibrium.
- At the start of a reaction, the reaction is under kinetic control, as the major product will be the one that is produced more quickly as a result of its lower activation energy.
- Given enough time, the reaction will be under thermodynamic control, as the dominant product will be the thermodynamically more stable product as a result of its lower free energy value.
